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Lecture.7 3.1 Chemical equations and balancing Chemical Equations : A chemical equation shows the reacting substances, called reactants, to the left of an arrow that points to the newly formed substances, called products: Reactants products Phases are often shown (s) for solid, (l),for liquid and (g) for gas. Compounds dissolved­ in water are designated (aq) for aqueous solution. Lastly numbers are place in front of the reactants or products to show the ratio in which they either combine or form. These numbers are called coefficients, and they represent numbers of individual atoms and molecules.

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C(s) + O 2 (g) CO 2 (g) The law of mass conservation states that matter is neither created nor destroyed, during a chemical reaction. This means that no atoms are lost or gained during any reaction. The chemical equation must therefore be balanced. In a balanced equation, each atom must appear on both sides of the arrow the same number of times. 2H 2 (g) + O 2 (g) 2H 2 O(g) (balanced )

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Mass Conservation & Balancing One of the important principles of chemistry is the law of mass conservation: matter (mass + energy) is neither created nor destroyed during a chemical reaction. The atoms present at the beginning of a reaction merely rearrange to form new molecules. This means that no atoms are lost or gained during a reaction. Therefore, the chemical equation must be balanced which means that each atom in the equation must appear the same number of times among the products as among the reactants. The equation for the formation of water is balanced: there are four hydrogen and two oxygen atoms before and after the arrow.

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A coefficient in front of a chemical formula tells us the number of times that element or compound must be counted. For example, 2H 2 O indicates two water molecules, which contain a total of four hydrogen and two oxygen atoms. By convention, the coefficient 1 is omitted so that the chemical equation above is typically written as: 2 H 2 (g) + O 2 (g) 2 H 2 O (g) (balanced) Additional Examples: C (s) + O 2 (g) CO 2 (g) (balanced) 3 O 2 (g) 2 O 3 (g) (balanced)

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: 3.2 Chemical Bonding Electron Shells We learned that an atom consists of a positively charged nucleus surrounded by moving, negatively charged electrons. According to the “shell model” of the atom, electrons behave as if they were arranged in concentric shells) around the nucleus). The adjacent figures are cutaway and cross-sectional views of shells (or energy levels!&) around the nucleus, resembling the Bohr’s model. There are seven shells available to the electrons in any atom, each with a limited capacity) for the number of electrons it can hold. From the first to the seventh shell, this capacity (consecutively is: 2, 8, 8, 18, 18, 32,and 32 electrons.

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:Electron-Dot Structures Based on the above atomic model, we learn the following points: 1. Electrons occupy the innermost shells first, where they are closest to the nucleus and possess minimum potential energy. Outer shells only get filled once the inner shells have reached their capacity for electrons. 2. Each shell corresponds to one period of the periodic table, and the capacity of a shell equals the number of elements in the corresponding period. 3. Each of the known noble gases (group 18 of the periodic table) has a filled shell of electron starting with Helium (2), Neon (2+8), argon(2+8+8), and so on. 4. Electrons in the outermost occupied shell are called valence electrons (from Latin valentia = strength), and they are said to occupy the atom’s valence shell.

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5. Valence electrons can be conveniently represented as a series of dots surrounding an atomic symbol. This notation is called the electron- dot structure. The adjacent figure shows the electron dot structures for the atoms needed in our forthcoming discussion of the ionic and covalent bonds.

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From the electron-dot structure of an atom, we can readily learn two important things relating to its bonding behavior. (a) Number of its valence electrons. (b) Number of paired valence electrons. Chlorine, for example, has three sets of paired electrons and one unpaired electron, and carbon has four unpaired electrons.

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Role of Valence Electrons in Bonding Unpaired Electrons Paired valence electrons are relatively stable.They resist change, and usually do not form chemical bonds with other atoms. For this reason, electron pairs in an electron-dot structure are called nonbonding pairs. By contrast, unpaired valence electrons have a strong tendency to participate in chemical bonding. By doing so, they become paired with an electron from another atom. The most stable electron arrangement for an atom is reached when all its valence electrons are paired so that its outermost occupied shell is filled to capacity An atom can fill its partially filled valence shell through bonding, via one of two methods: a) Sharing valence electrons with another atom. b) Transferring valence electrons to another atom. We will see that this leads to three types of chemical bonds: ionic, covalent, and metallic.

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The octet rule Atoms tend to form chemical bonds with each other, so that each has eight electrons in its valence shell, similar to the electron configuration of a noble gas. Example: Sodium, being a group 1 element, has one valence electron. If an atom has only one or only a few electrons in its valence shell, it will tend to lose its outer-shell electrons so that the next shell inward, which is filled, becomes the outermost occupied shell. Then, the atom will have a filled valence shell. Sodium readily gives up the single electron in its third shell. This makes the second shell, which is already filled to capacity, sodium’s outermost occupied she ll

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3.3 Ion Formation and Bonding Forming Ions We can use the shell model to deduce the type of ion an atom tends to form. According to this model, atoms tend to lose or gain electrons so as to form an outermost occupied shell that is filled to capacity. We can use the periodic table as a quick reference when determining the type of ion an atom tends to form. As shown in the figure, a group-1 atom, for example, has only one valence electron, and so tends to form a + 1 ion. A group-17 atom has room for one additional electron in its valence shell,therefore tending to form a -1 ion. Atoms of the noble gases tend not to form any type of ion because their valence shells are already filled to capacity.

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The above figure indicates that the attraction of an atom’s nucleus for its valence electrons is weakest in elements on the left in the periodic table and strongest in elements on the right.

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The above figure further indicates that positive ions generally derive from metallic elements (on the left side of the table), and negative ions from nonmetallic elements (on the right side ). The shell-model explanation of how ions form works well for groups 1 and 2, and groups 13 through 18. This model, however, is too simplistic to work well for the transition metals (groups 3-12) or the inner transition metals. In general, the atoms of these metals tend to form positive ions, but the number of electrons lost varies. Example: Depending on conditions, an iron atom may lose two electrons (forming Fe 2+ ) or three electrons (forming Fe 3+ ). T he nucleus of a noble gas atom pulls so strongly on its valence electrons that they are very difficult to remove. Because there is no room left in the valence shell of a noble-gas atom, no additional electrons are gained. Thus, a noble gas atom tends not to form ions of any sort

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We saw that atoms form ions by losing or gaining electrons. We should note now that molecules can also form ions. This usually happens when a molecule loses or gains a proton, which is the same thing as one hydrogen ion (H + ). Examples: a) A water molecule (H 2 O) can gain a hydrogen ion (H + ) to form the hydronium ion (H 3 O + ) :

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b ) A carbonic acid molecule (H 2 CO 3 ) can lose two protons to form the ( carbonate ion (CO 3 -2 Hydronium and carbonate ions are examples of polyatomic ions, which are molecules that carry a net electric charge. These and other common polyatomic ions are included in the following table :

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The Ionic Bond When an atom that tends to lose electrons comes in contact with an atom that tends to gain them, the result is an electron transfer and the formation of two oppositely charged ions. This is what happens when sodium and chlorine are combined. The two oppositely charged ions are attracted to each other by the electric force, which holds them together in what is called an ionic bond, resulting in the chemical compound sodium chloride. Any chemical compound containing ions is referred to as ionic compound. All ionic compounds are completely different from the elements from which they are made.

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Examples: a) In sodium chloride, there is one sodium (1+) ion for every chloride (-1) ion. This results in the formula NaCl. b) In calcium fluoride, one calcium (+2) ion must take on two fluoride (-1) ions. This results in the formula CaF 2. c) In aluminum oxide Al 2 O 3, one aluminum (+3) ion must take on 1½ oxide (-2) ions. In other words, two aluminum ions need 3 oxide ions (6-) to balances their total 6+ charge, resulting in the formula Al 2 O 3. This is the mineral that composes rubies and sapphires

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An ionic compound typically contains a multitude of ions grouped together in a highly Ordered three dimensional array). In sodium chloride, for example, each sodium ion is surrounded by six chloride ions and each chloride ion is surrounded by six sodium ions. Overall, there is one sodium ion for each chloride ion, but (there are no identifiable sodium-chloride pairs. Such an orderly array of ions is known as an ionic crystal

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On the atomic level, the crystalline structure of sodium chloride is cubic, which is why macroscopic crystals of table salt are also cubic. Smash a large cubic sodium chloride crystal with a hammer, and what do you get? Smaller cubic sodium chloride crystal

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3.4 The Covalent Bond This type of electrical attraction in which atoms are held together by their mutual attraction for shared electrons is called a covalent bond where “co-” signifies sharing and “-valent” refers to the fact that valence electrons are being shared. A compound composed of atoms held together by covalent bonds is a covalent compound. The fundamental unit of most covalent compounds is a molecule, which we can now formally define as any group of atoms held together by covalent bonds. Atoms held together with a covalent bond normally have two types of electron pairs:

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a) Bonding pair; this refers to any pair arising from the formation of a covalent bond. This means that the two electrons come from two different atoms. b) Nonbonding pair; this refers to other pairs existing in the electron-dot structure of an individual atom. Therefore, both electrons in this pair belong to the same atom. The electron-dot structures for covalent compounds are often expressed as a straight line representing the bonding pair of electrons..

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In some cases where the nonbonding electron pairs play no significant role in the process at hand, these electrons are left out of the symbol. Example: In both cases, the straight line represents the two shared electrons

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A covalent bond, by contrast, is formed when two atoms that tend to gain electrons are brought into contact with each other. Atoms that tend to form covalent bonds are, therefore, primarily atoms of the nonmetallic elements in the upper right corner of the periodic table (with the exception of the noble gas elements. Hydrogen tends to form covalent bonds because, unlike the other group 1 elements, it has a fairly strong attraction for an additional electron, and two hydrogen atoms covalently bond to form a hydrogen molecule, H 2, as shown.

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The number of covalent bonds an atom can form is equal to the number of additional electrons it can attract, which is the number it needs to fill its valence shell. Examples: a) Hydrogen attracts only one additional electron, and so it forms only one covalent bond. b) Oxygen, which attracts two additional electrons, finds them when it encounters two hydrogen atoms and bonds with them covalently to form water, H 2 O, as shown. With this bonding, the oxygen atom has access to two additional electrons, while each hydrogen atom has access to one additional electron. Thus, each atom achieves a filled valence shell

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d) A carbon atom can attract four additional electrons, and is thus able to form four covalent bonds, as occurs in methane, CH 4. It is possible to have more than two electrons shared between two atoms. Examples: a) Molecular oxygen (O 2 ) consists of two oxygen atoms connected by four shared electrons. This arrangement is called a double covalent bond or, for short, a double bond..

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b) The covalent compound carbon dioxide (CO 2 ) consists of two double bonds connecting two oxygen atoms to a central carbon atom. c) Some atoms can form triple covalent bonds, in which six electrons (three from each atom) are shared. One example is molecular nitrogen (N 2 ). Multiple covalent bonds higher than triple are not commonly observed

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3.5 Polar Bonds and Polar Molecules In a covalent bond, if the two atoms are identical, their nuclei have the same positive charge and, therefore, the electrons are shared evenly We can represent these electrons as being centrally located by using an electron-dot structure in which the electrons are situated exactly halfway between the two atomic symbols. Example: hydrogen, H : H. In a covalent bond between non-identical atoms, the nuclear charges are different and, consequently, the bonding electrons may be shared unevenly..

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Example: In a hydrogen-fluorine bond, electrons are more attracted to fluorine’s greater nuclear charge. In the electron-dot structure, this is represented by showing the two electrons closer to the fluorine atom (H : F). The bonding electrons spend more time around the fluorine atom. For this reason, the fluorine side of the bond is slightly negative and, consequently, the hydrogen side of the bond is slightly positive

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This separation of charge is called a dipole, and is represented either by the characters E- and E+, read “slightly negative” and “slightly positive,” respectively, or by a crossed arrow pointing to the negative side of the bond: The strength of an atom’s ability to pull bonding electrons is called electronegativity, and is an experimentally measured quantity, ranging from 0.7 to 3.98 (see next figure)..

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Thus, in hydrogen fluoride, fluorine has a greater electronegativity, or pulling power, than hydrogen. Electronegativity is greatest for elements at the upper right of the periodic table and lowest for elements at the lower left. Noble gases are not considered because they (generally) are inert – their outermost shells being already filled, and the electrons in those shells being tightly held

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When the two atoms in a covalent bond have the same electronegativity (as is the case with H 2 ), no dipole is formed, and the bond is classified as nonpolar. When the electronegativities of the atoms differ, a dipole may form (as with HF), and the bond is classified as a polar. The polarity of a bond is stronger with a greater difference between the electronegativities of the bonded atoms.

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It is important to note that there is no clear-cut distinction between ionic and covalent bonds. Rather, there is a gradual change from one to the other as the atoms that bond are located farther and farther apart in the periodic table. This continuum is illustrated in the figure, where the bonds are shown in order of decreasing polarity from left to right with the size of the crossed arrow representing the degree of polarity..

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Atoms on opposite sides of the periodic table have great differences in electronegativity, and hence the bonds between them are highly polar (i.e., ionic). Nonmetallic atoms of the same type have the same electronegativities, and so their bonds are nonpolar covalent. The polar covalent bond, with its uneven sharing of electrons and slightly charged atoms, is between these two extremes